Chemistry Reference
In-Depth Information
H
2
C
CH
2
H
2
N
CH
2
CH
2
CH
2
N
H
2
H
2
N
CH
2
CH
2
N
HCH
3
H
2
N
CH
2
N
HCH
2
CH
3
H
N
N
H
CH
2
H
2
C
H
2
C
N
H
H
2
C
H
2
C
CH
2
H
2
C
2
H
2
C
CH
2
CH
2
N
H
CH
3
H
2
N
N
H
2
H
2
N
N
H
H
2
N
N
H
CH
3
M
M
M
M
Figure 4.2
Chelation modes for simple linear diamines of common empirical formula H
10
C
3
N
2
. Also included
for contrast is a related molecule where the linear diamine to its immediate left has the five atoms in
the chain confined in an organic ring; it can adopt only monodentate coordination to the metal.
Regardless of the identification of the effects, one obvious outcome is that the 'octahedral'
[CoBr(NH
3
)
5
]
2
+
ion is not an ideal octahedron.
If we replace the cobalt(III) centre in [Co(NH
3
)
6
]
3
+
by a different metal ion such as
Rh(III) or Ni(II), what we discover is that the octahedral shape is retained, but the average
M N distance changes. Obviously, each metal has a preferred distance between its centre
and the donor atom. We can understand this in terms of both the size of the metal ion
- the larger, the longer the bonds - and in terms of the charge on the metal ion. With
variation in charge (or oxidation state) two factors operate; first, the number of d electrons
differ, which can influence both shape preference and repulsive terms related to the number
of valence shell electrons; second, the charge changes, influencing simple electrostatic
interactions. The effects are exemplified by examining one metal in two oxidation states;
for example, Co(III) N bonds are invariably shorter than Co(II) N bonds. Because a
coordinate bond links two different atomic centres, it is reasonable to expect that preferred
metal-donor distance changes with donor also. We have seen already how a Co(III) N
bond distance differs from a Co(III) Br
−
bond distance. This is a universal observation -
the preferred metal-donor distance varies with the type of donor even when the metal is
fixed. In fact, the effect is quite subtle, as the Co(III) NH
3
distance differs from that found
for Co(III) NH
2
CH
3
, even though they both form a Co N bond. This can be accounted
for by two factors; first, ammonia is a smaller, less bulky molecule than methylamine;
second, the basicity of ammonia and methylamine differ, affecting their relative capacities
to act as a lone pair donor (Lewis base) in forming a coordinate covalent bond to the metal
ion (Lewis acid). The size and also the shape of ligands can force much more dramatic
effects upon their complexes, as we shall see, so it is apparent that ligand geometry and
rigidity are important. We can illustrate this in part for three simple diamines, all of formula
H
10
C
3
N
2
. Chelation of each of these in turn would produce a six, five and four-membered
ring (Figure 4.2). With five-membered chelate rings usually being more stable than six- or
four-membered rings, it is not surprising to find that shapes of complexes with the ligand
forming a five-membered chelate ring are less distorted than those with the other ligands.
Further, if we join the terminal carbon and nitrogen atoms in the latter example to form a
small five-membered heterocyclic ring (Figure 4.2), the ligand shape is now such that only
one of the two amine groups can bind to a single metal centre at any one time, the other
having its lone pair directed in an inappropriate direction for chelation.
